In simple terms
A friendly intro before the formal notes — no formulas yet.
The Proton Dance
Acids and bases are defined by their ability to donate or accept protons, respectively. This simple exchange governs everything from the acidity of a solution (pH) to the function of biological buffers.
Imagine a ballroom where dancers represent molecules and a special hat represents a proton (H⁺). An 'acid' dancer is one who gives their hat to another dancer. A 'base' dancer is one who accepts a hat. A buffer solution is like a large group of dancers who are very good at quickly passing the hat back and forth amongst themselves, ensuring that the number of dancers left without a hat at any one time remains almost constant, even if new dancers join or leave the floor.
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Brønsted–Lowry: acid proton donor, base proton acceptor.
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Ka, pKa, pH — weak acid partial dissociation.
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Buffer resists pH change — weak acid + conjugate base.
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Titration curves — equivalence point pH depends on acid/base strength.
Explore the concept
Use the live diagram, PhET or GeoGebra sim, and synced steps — play it, drag controls, or tap a step.
Step-synced diagram — highlights what to look for in the simulation above.
Brønsted–Lowry: acid proton donor, base…
Brønsted–Lowry: acid proton donor, base proton acceptor.
2 more simulations for this topic — run them in the Simulations section below
Simulations
Every simulation here runs the real model — try the steps on a card, then check what you see against the notes.
2 simulations
- PhETCore9701 25.1
pH Scale
Pick a liquid, dilute it, and see pH alongside the H₃O⁺ and OH⁻ concentrations.
Try this
- On Micro choose battery acid, read [H₃O⁺] and [OH⁻] and multiply them together.
- Drain to 0.1 L, top up with water to 1.0 L and note the change in pH.
- On My Solution set pH 7 and compare the two concentrations.
Look for [H₃O⁺][OH⁻] stays at 1.0 × 10⁻¹⁴ mol² dm⁻⁶ (K_w at 25 °C) whatever the pH, and each pH unit is a tenfold change in [H⁺].
Simulation by PhET Interactive Simulations, University of Colorado Boulder · Licensed to MarkScheme (public licence CC BY-NC 4.0 since 2026-03-30)
- GeoGebraCore9701 25.1
Weak acid–strong base titration curve
A pH curve for a weak acid titrated with a strong base, with sliders for the acid volume, both concentrations and Ka.
Try this
- Find the equivalence point, where the curve is steepest.
- Raise Ka (a stronger acid) and watch the start of the curve drop.
- Move the pH marker to half the equivalence volume and read the pH.
Look for The equivalence point lies above pH 7, and at half-neutralisation pH = pKa — the flat buffer region.
Sune Hvidtfeldt Håkansson · GeoGebra · GeoGebra Terms of Service
Key formulas
Tap any symbol to reveal exactly what it means and its units.
$pH = -log_{10}[H^+(aq)]$
$For a weak acid HA: HA(aq) \rightleftharpoons H^+(aq) + A^-(aq) \ K_a = \frac{[H^+(aq)][A^-(aq)]}{[HA(aq)]}$
$pH = pK_a + log_{10}\left(\frac{[A^-]}{[HA]}\right)$
Full topic notes
Formal explanation with the rigour you need for the exam.
1. The Brønsted–Lowry Theory
The Brønsted–Lowry theory provides a more general definition of acids and bases than the Arrhenius theory. It focuses on the transfer of a proton, which is simply a hydrogen ion, H⁺. An acid is a proton donor, and a base is a proton acceptor. When an acid donates a proton, the species that remains is its conjugate base. When a base accepts a proton, it forms its conjugate acid. Together, they form a conjugate acid-base pair.
Consider the reaction of ammonia with water: . Here, water donates a proton to ammonia. Therefore, acts as an acid and acts as a base. The ammonium ion, , is the conjugate acid of , and the hydroxide ion, , is the conjugate base of . Substances like water that can act as both an acid and a base are called amphoteric.
2. pH, Ka, and pKa: Quantifying Acidity
The pH scale provides a convenient way to express the concentration of hydrogen ions in a solution. It is a logarithmic scale, meaning a change of one pH unit represents a tenfold change in .
pH = -log_{10}[H^+(aq)]
Strong acids, like HCl, fully dissociate in water, so for a $0.1 \ mol \ dm^{-3}$ solution of HCl, $[H^+] = 0.1 \ mol \ dm^{-3}$. Weak acids, like ethanoic acid (), only partially dissociate. This partial dissociation is an equilibrium, described by the acid dissociation constant, Ka.
A larger Ka value indicates a greater extent of dissociation and therefore a stronger acid.
pKa is used for convenience: $pK_a = -log_{10}(K_a)$.
A smaller pKa value indicates a stronger acid.
The ionic product of water, $K_w = [H^+][OH^-] = 1.0 \times 10^{-14} \ mol^2 \ dm^{-6}$ at 298 K, links the concentrations of H⁺ and OH⁻ in any aqueous solution.
3. Buffer Solutions
A buffer solution is a remarkable chemical system that resists changes in pH when small quantities of acid or alkali are added. Acidic buffers are typically made from a weak acid and its conjugate base (usually from a salt). For example, a mixture of ethanoic acid () and sodium ethanoate (). The solution contains a large reservoir of both the weak acid and its conjugate base.
$pH = pK_a + log_{10}\left(\frac{[A^-]}{[HA]}\right)$
Adding acid (H⁺): The excess H⁺ ions are removed by reacting with the conjugate base: . The equilibrium shifts to the left.
Adding alkali (OH⁻): The added OH⁻ ions are neutralised by the weak acid: . The equilibrium shifts to the right.
In both cases, the large reservoirs of HA and A⁻ ensure the change in the ratio is small, thus the pH change is minimal.
The pH of a buffer is determined by the pKa of the weak acid and the ratio of the concentrations of the conjugate base and weak acid.
When asked to explain how a buffer works, you must write two separate equations: one showing the reaction with added H⁺ and one showing the reaction with added OH⁻. State that there are large reservoirs of both the weak acid and its conjugate base.
4. Acid-Base Titration Curves
A titration curve is a graph of pH against the volume of titrant added. The shape of the curve provides valuable information about the strengths of the acid and base involved. Key features include the initial pH, the buffer region (for weak species), the equivalence point (steepest part of the curve), and the final pH.
Strong Acid - Strong Base: Equivalence point at pH 7. Steep vertical section.
Weak Acid - Strong Base: Initial pH is higher (weak acid). A buffer region exists before the equivalence point. Equivalence point is above pH 7 due to the hydrolysis of the conjugate base formed ().
Strong Acid - Weak Base: Equivalence point is below pH 7 due to the hydrolysis of the conjugate acid formed ().
Indicator Choice: A suitable indicator must have a colour change range (pKin) that falls entirely within the steep vertical section of the titration curve.
When sketching titration curves, always label your axes (pH on y-axis, Volume of titrant/cm³ on x-axis). Mark the volume and pH at the equivalence point. For weak acid/base titrations, also indicate the approximate pH at the half-equivalence point, where .
Worked examples
See the formulas applied — reveal one step at a time, like the exam.
Calculate the pH of a $0.0500 \ mol \ dm^{-3}$ solution of propanoic acid at 298 K. The for propanoic acid is $1.35 \times 10^{-5} \ mol \ dm^{-3}$.
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Write the Ka expression:
A buffer solution is made by dissolving 12.3 g of sodium ethanoate () in 250 cm³ of $0.800 \ mol \ dm^{-3}$ ethanoic acid. Calculate the pH of the resulting buffer solution. ( of ; for $CH_3COOH = 1.75 \times 10^{-5} \ mol \ dm^{-3}$)
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Calculate moles of conjugate base (ethanoate):
How it all connects
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Glossary
Key terms for this topic — skim now; the Check step will test them.
- Brønsted–Lowry acid
A proton (H⁺) donor.
- Brønsted–Lowry base
A proton (H⁺) acceptor.
- Conjugate acid-base pair
A pair of species that differ by a single proton (H⁺). For example, HA and A⁻.
- PH
The negative logarithm to the base 10 of the hydrogen ion concentration. $pH = -log_{10}[H^+(aq)]$. A lower pH indicates higher acidity.
- Acid dissociation constant, Ka
An equilibrium constant for the dissociation of a weak acid in water. For , . A larger Ka indicates a stronger weak acid.
- PKa
The negative logarithm to the base 10 of the acid dissociation constant, Ka. $pK_a = -log_{10}(K_a)$. A smaller pKa indicates a stronger acid.
- ionic product of water,
The equilibrium constant for the auto-ionisation of water. . At 298 K, its value is $1.0 \times 10^{-14} \ mol^2 \ dm^{-6}$.
- Buffer solution
A solution that resists changes in pH upon the addition of small amounts of acid or alkali. It consists of a weak acid and its conjugate base, or a weak base and its conjugate acid.
- Henderson-Hasselbalch equation
An equation used to calculate the pH of a buffer solution: $pH = pK_a + log_{10}(\frac{[A^-]}{[HA]})$. It is not explicitly required in the syllabus but is a very useful tool.
Name it
Read the meaning, then pick which of this lesson’s terms it describes. Miss one and you see what your choice really means.
A pair of species that differ by a single proton (H⁺). For example, HA and A⁻.
Quick check
Write your answer first, then compare it with the model one — the gap is what you would have lost.
Teach it back
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Teach it back
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Revision flashcards
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Key takeaways
Review these before you close the topic — retrieval beats re-reading.
A larger Ka value indicates a greater extent of dissociation and therefore a stronger acid.
pKa is used for convenience: $pK_a = -log_{10}(K_a)$.
A smaller pKa value indicates a stronger acid.
The ionic product of water, $K_w = [H^+][OH^-] = 1.0 \times 10^{-14} \ mol^2 \ dm^{-6}$ at 298 K, links the concentrations of H⁺ and OH⁻ in any aqueous solution.
Practice — then mark it
The whole point: a real Cambridge question, marked mark-by-mark.
Ammonium carbonate undergoes an acid-base reaction with NaOH(aq). Explain this statement.
Buffer solutions are used to regulate pH. Write two equations to describe how a solution containing HC2O4¯ ions acts as a buffer solution when small amounts of acid or alkali are added.
Extra simulations & links
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Frequently asked
Checkpoint
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